Question

Arrange the following groups of atoms in order of increasing size. a. \(\mathrm{Rb}, \mathrm{Na}, \mathrm{Be}\) b. \(\mathrm{Sr}, \mathrm{Se}, \mathrm{Ne}\) c. \(\mathrm{Fe}, \mathrm{P}, \mathrm{O}\)

Short answer

The order of increasing atomic size for the given groups of atoms are: a. \(\mathrm{Be} < \mathrm{Na} < \mathrm{Rb}\) b. \(\mathrm{Ne} < \mathrm{Se} < \mathrm{Sr}\) c. \(\mathrm{O} < \mathrm{P} < \mathrm{Fe}\)

Step-by-step solution

Locate elements in the periodic table

Using a periodic table, determine the position of each element: - Rb (Rubidium) belongs to period 5 and group 1 (alkali metals) - Na (Sodium) belongs to period 3 and group 1 (alkali metals) - Be (Beryllium) belongs to period 2 and group 2 (alkaline-earth metals)

Order elements by size

Based on the position of these elements in the periodic table, we observe: - Rb has the greatest size (moving down the group) - Na is larger than Be (moving to group 1, and Na is below Be) Thus, the order of increasing atomic size is: \(\mathrm{Be} < \mathrm{Na} < \mathrm{Rb}\) b. Arrange \(\mathrm{Sr}, \mathrm{Se}, \mathrm{Ne}\) in order of increasing size.

Locate elements in the periodic table

Identify the positions of each element in the periodic table: - Sr (Strontium) belongs to period 5 and group 2 (alkaline-earth metals) - Se (Selenium) belongs to period 4 and group 16 (chalcogens) - Ne (Neon) belongs to period 2 and group 18 (noble gases)

Order elements by size

Considering their positions in the periodic table, we find: - Sr has the largest size (moving down the group) - Se is larger than Ne (both in the same period, Se in a group further to the left) Therefore, the order of increasing atomic size is: \(\mathrm{Ne} < \mathrm{Se} < \mathrm{Sr}\) c. Arrange \(\mathrm{Fe}, \mathrm{P}, \mathrm{O}\) in order of increasing size.

Locate elements in the periodic table

Locate the elements in the periodic table and note their positions: - Fe (Iron) belongs to period 4 and group 8 (transition metals) - P (Phosphorus) belongs to period 3 and group 15 (pnictogens) - O (Oxygen) belongs to period 2 and group 16 (chalcogens)

Order elements by size

Based on the relative position of these elements in the periodic table, we find: - Fe has the greatest size (furthest down the table) - P is larger than O (both in the same group, P is further down) Hence, the order of increasing atomic size is: \(\mathrm{O} < \mathrm{P} < \mathrm{Fe}\)

Key concepts

Atomic Size

Atomic size refers to the volume of space an atom occupies, often described by the term "atomic radius." The atomic size changes based on an element's position in the periodic table. Generally, as you move down a group in the periodic table, atomic size increases. This is because additional electron shells are added, which means the electrons are further from the nucleus.

• In the exercise, for example, Rubidium (Rb), Sodium (Na), and Beryllium (Be) were arranged by atomic size. Rb is in period 5, Na in period 3, and Be in period 2. Consequently, Rb has the largest atomic size due to more electron shells compared to Na and Be.
• Similarly, for Strontium (Sr), Selenium (Se), and Neon (Ne), Sr has the largest atomic size because it sits in a lower period.
• In the last group with Iron (Fe), Phosphorus (P), and Oxygen (O), Fe has the largest atomic size, again due to being further down the table.

Recognizing these patterns in atomic size helps in predicting the behavior of an element, such as its reactivity and bonding with other elements. Understanding trends in atomic size across the periodic table is vital for understanding how specific elements will behave.

Periodic Trends

Periodic trends are systematic patterns observed in the periodic table that aid in predicting the behavior and properties of elements. Trends include changes in atomic size, electronegativity, ionization energy, and electron affinity as you move across a period or down a group.

• Atomic size decreases from left to right across a period. This occurs because, although more protons and electrons are added, the increased positive charge pulls electrons closer to the nucleus, with no extra electron shells in the same period.
• Conversely, as discussed earlier, atomic size increases as you move down a group. From our exercise, the order of elements was based on such observations: Be < Na < Rb, Ne < Se < Sr, and O < P < Fe.
• These trends are crucial for anticipating how different elements interact and bond with each other.

Understanding periodic trends allows chemists to make accurate predictions about the properties of new and existing elements, facilitating the study of chemical reactions and the creation of new compounds. Grasping these trends equips students to relate elemental behavior to their positioning in the periodic table.

Element Groups

Element groups refer to columns in the periodic table and contain elements with similar physical and chemical properties. Elements in the same group have the same number of electrons in their outermost shell, leading to similar ways in which they interact or bond with other elements.

• For instance, in the exercise, Rubidium (Rb) and Sodium (Na) both belong to the alkali metals group. They have one electron in their outer shell, making them highly reactive.
• Strontium (Sr) and Beryllium (Be) are found in the alkaline-earth metals group and have two electrons in the outer shell, affecting how they form compounds.
• Neon (Ne) is a noble gas, known for its full outer electron shell, making it very stable and unreactive.

These groups help predict how elements may behave or react under certain conditions, providing a roadmap for understanding chemical interactions. Element groups are a fundamental concept in chemistry, allowing scientists to grasp the connectivity and reactivity of different metals, nonmetals, and metalloids.